Surface Structural Changes of Perovskite Oxides during
Oxygen Evolution in Alkaline Electrolyte
MCHIVEs
ASSACHUSETV1S INS71TrE~
F TECHNOLoY
by
UN 2 5 2
Kevin J. May
BRA S
B.Eng., Engineering Physics, McMaster University (2010)
SUBMITTED TO THE DEPARTMENT OF MECHANICAL ENGINEERING
IN PARTIAL FULFILLMENT OF THE REQUIREMENTS FOR THE DEGREE OF
MASTER OF SCIENCE IN MECHANICAL ENGINEERING
at the
MASSACHUSETTS INSTITUTE OF TECHNOLOGY
June 2013
@
2013 Massachusetts Institute of Technology
All rights reserved.
..
.......................
...
Department of Mechanical Engineering
May 10, 2013
Signature of Author.......... .............
.....
C ertified by...........................................
Gail
Accepted by.........
.
7.. .Yang
Shao-Horn
Kendall Professor of Mechanical Engineering
Thesis Supervisor
...................
....... ........................
David E. Hardt
Ralph E. and Eloise F. Cross Professor of Mechanical Engineering
Chairman, Department Graduate Committee
2
Surface Structural Changes of Perovskite Oxides during
Oxygen Evolution in Alkaline Electrolyte
by
Kevin J. May
Submitted to the Department of Mechanical Engineering
on May 10, 2013 in Partial Fulfillment of the
Requirements for the Degree of Master of Science in
Mechanical Engineering
Abstract
(BSCF82) are among the most active
catalysts for the oxygen evolution reaction (OER) in alkaline solution reported to date.
In this work it is shown via high resolution transmission electron microscopy (HRTEM)
and Raman spectroscopy that oxides such as BSCF82 rapidly undergoes amorphization
at its surface under OER conditions, which occurs simultaneously with an increase in
the pseudocapacitive current and OER activity. This amorphization was not detected at
potentials below those where significant OER current was observed. Lower
concentrations of Sr'- and Ba' are found in the amorphous regions of BSCF82.
Perovskite oxides with lower OER activities such as LaCoO 3 (LCO) and LaMnO 3
(LMO) remained crystalline under identical electrochemical conditions. In addition, the
OER activity and tendency for amorphization are found to correlate with the oxygen pband center as calculated using density functional theory. This work illustrates that the
surface structure and stoichiometry of oxide catalysts can differ significantly from the
bulk during catalysis, and that understanding these phenomena is critical for designing
highly active and stable catalysts for the OER.
Perovskite oxides such Bao.5 Sr 0 Co0 8 FeO 803
6
Thesis Supervisor: Yang Shao-Horn
Title: Gail E. Kendall Professor of Mechanical Engineering
3
4
List of Publications
1. Grimaud, A., Carlton, C.E., Risch, M., Hong, W.T., May, K.J. & Shao-Horn, Y.
On the Influence of Cobalt and Manganese Coordination on the Catalytic Activity of
Perovskites for Oxygen Evolution. Submitted.
2. Grimaud, A., May, K.J., Carlton, C.E., Lee, Y.-L., Risch, M., Zhou, J. & ShaoHorn, Y. Double Perovskites as a New Family of Highly Active Catalysts For
Oxygen Evolution in Alkaline Solution. Submitted.
3. Ming, T., Suntivich, J., May, K.J., Stoerzinger, K.A. & Shao-Horn, Y. Visible
Light Photo-Oxidation in Au Nanoparticle Sensitized SrTiO,:Nb Photoanode.
Submitted.
4. Risch, M., Grimaud, A., May, K.J., Stoerzinger, K.A., Chen, T.J., Mansour, A.N.,
& Shao-Horn, Y. Structural Changes of Cobalt-based Perovskites upon Water
Oxidation Investigated by EXAFS. The Journal of Physical Chemistry C 117, 86288635 (2013).
5. May, K. J., Carlton, C.E., Stoerzinger, K.A., Risch, M., Suntivich, J., Lee, Y.-L.,
Grimaud, A., & Shao-Horn, Y. The Influence of Oxygen Evolution during Water
Oxidation on the Surface of Perovskite Oxide Catalysts. The Journal of Physical
Chemistry Letters 3, 3264-3270 (2012).
6. Lee, Y., Suntivich, J., May, K. J., Perry, E. E. & Shao-Horn, Y. Synthesis and
Activities of Rutile Ir0 2 and RuO 2 Nanoparticles for Oxygen Evolution in Acid and
Alkaline Solutions. The Journal of Physical Chemistry Letters 3, 399-404 (2012).
7. Suntivich, J., May, K. J., Gasteiger, H. A., Goodenough, J. B. & Shao-Horn, Y. A
Perovskite Oxide Optimized for Oxygen Evolution Catalysis from Molecular Orbital
Principles. Science 334, 1383-1385 (2011).
5
6
Acknowledgements
I would first like to acknowledge my advisor, Professor Yang Shao-Horn, for not
only her academic and research-oriented guidance but also for giving unwavering
support in all other aspects of my graduate career. You have believed in my abilities as
a student and researcher, and fostered my independence and creativity; I have learned
so much more than just science and research during my time at MIT because of this.
Thank you.
Funding for this work is gratefully acknowledged. Support was received from the
MRSEC Program of the National Science Foundation (DMR-0819762), Eni S.p.A. under
the Eni-MIT Alliance Solar Frontiers, and the DOE Hydrogen Initiative Program (DEFG02-05ER15728). I also received support from the Natural Sciences and Engineering
Research Council of Canada (PGS-M).
I would like to also extend my appreciation and thanks to my fellow lab members
in the Electrochemical Energy Lab for their friendship and support. Both present and
past, the students and postdoctoral researchers in this group are the best I could ever
imagine working with. I appreciate the support of everyone in the group, and would like
to mention some members in particular who have contributed to the scientific work in
this thesis: Jin Suntivich, for his enthusiasm in teaching the fundamentals of
electrochemistry and for providing mentorship when I was beginning my graduate
studies, which significantly shaped my way of scientific thinking and how I approach
research; Chris Carlton, for his TEM expertise and helpful discussions in the broader
subject of materials science; Kelsey Stoerzinger for her Raman spectroscopy
contributions; Marcel Risch for assisting in electrode preparation for Raman
spectroscopy and for his expertise in amorphous cobalt oxides and electrochemistry;
Yueh-Lin Lee for density functional theory calculations, and Alexis Grimaud for alwayshelpful discussions in solid state chemistry and synthesis of perovskite oxide compounds.
I also thank my family for their love and support, for making me the person I am
today and also for giving me all the opportunities I needed to succeed. My parents never
faltered in fostering my curiosity and encouraging me unconditionally in my endeavors.
To Megan and Brandon, I'm lucky to have you as siblings, and I know I can always rely
on you for support.
Finally, I want to give my love and gratitude to Sana, for being my other half
through thick and thin. You have given me the strength and confidence to persevere
through any challenges that arise. This has helped me so much, perhaps even more than
you know. Thank you for being who you are and for being my partner in everything.
7
8
Table of Contents
A b st r a ct ........................................................................................................................
3
List of Publications .....................................................................................................
5
A cknow ledgem ents.................................................................................................
7
List of Figures .............................................................................................................
11
List of Tables ..............................................................................................................
15
Introduction ...........................................................................................................
17
1.1 M o tiv a tio n ........................................................................................................
17
1
2
3
1.2
The Oxygen Evolution Reaction: Historical Overview ......................................
18
1.3
Perovskite Oxides as Oxygen Evolution Catalysts........................................
23
1.4
Scope of this Thesis.......................................................................................
26
Experim ental..........................................................................................................29
2.1
Oxide Synthesis and Characterization...........................................................
29
2.2
Electrode Preparation ..................................................................................
31
2.3
Electrochem ical M easurements.......................................................................31
2.4
Characterization of Electrochem ically-Cycled Electrodes..............................
32
2.5
Density Functional Theory Calculations ......................................................
34
Results and D iscussion.....................................................................................
35
3.1
M aterial Characterization ............................................................................
35
3.2
Electrochem ical M easurem ents.......................................................................37
3.3
Transm ission Electron M icroscopy .................................................................
42
3.4
Raman Spectroscopy and Potential-Dependent Amorphization ....................
48
3.5
Am orphization and Pseudocapacitive Current ..............................................
51
3.6
0 p-band Center: A Descriptor from Density Functional Theory ..........
54
4
Conclusions and Looking Forward ................................................................
57
5
R e fe re n c e s ..............................................................................................................
61
9
10
List of Figures
Figure 1.1: Free energies of OER intermediates on an oxide surface, for the case of
zero applied potential and 1.23 V applied potential. The specific intermediates assumed
are not considered here. The "ideal" catalyst is shown in red, while a more realistic
catalyst is shown in blue. The ideal catalyst has a flat energy landscape at the
thermodynamic equilibrium potential of 1.23 V vs. RHE. Figure adapted from work by
R ossm eisl et al." and D au et al.'..................................... . .. ... ... .. .. ... .. ... .. ... .. ... .. ... .. . . . 23
Figure 1.2: Cubic perovskite structure. The central B-site transition metal (dark blue)
is octahedrally coordinated by oxygen (red), with the A-site alkaline earth or lanthanide
metal (yellow) at the corners in a 12-fold coordination by oxygen.............................24
Figure 1.3: Qualitative molecular orbital diagram for octahedrally-coordinated
transition metal ion. Asterisks denote antibonding states. In bold are the eg and t2g
states of interest in catalysis. Figure adapted from Cox.................... . .. ... ... .. ... .. ... . . . . 25
Figure 1.4: Volcano plot showing the relationship between OER catalytic activity,
defined as the overpotential at 50 iA cm- 2 (estimated true oxide surface area), and the
10
eg orbital occupancy of the B-site transition metal ion. Figure from previous work. ....26
Figure 3.1: X-ray diffraction data for the oxides studied in this work. LMO, LCO data
were gathered by Jin Suntivich. LSC46 data were gathered by Ethan J. Crumlin. SCF82
35
data were gathered by Alexis Grim aud. .....................................................................
Figure 3.2: Experimental XRD peak locations (red) with database peak positions for
LSC46 (gray) and Si (green). Si was included to give reference diffraction peaks. ......... 36
Figure 3.3: Cyclic voltammetry for (a) LMO, (b) LCO, (c) LSC46, (d) BSCF46, (e)
BSCF82 and (f) SCF82. Number labels n indicate the nth cycle of cyclic voltammetry
(1.1-1.7 V versus RHE, 10 mV/s). All measurements were performed in 0 2-saturated 0.1
M KOH, corrected for ohmic losses and normalized to the SEM-estimated surface area
of oxides, and used an oxide loading of 0.25 mg/cm2 supported on glassy carbon. Figure
38
adapted from published w ork.' ..................................................................................
and (b) LSC46; (c) Potentiostatic
Figure 3.4: Heavy cycling of (a) BSCF82
measurements on BSCF82, LSC46, LCO and LMO; (d) Tafel plot containing cyclic
voltammetry at the nth cycles (hollow markers) and potentiostatic data (solid markers)
for comparison. Figure adapted from published work."..................... .. ... .. ... ... .. .. ... .. . . . 39
Figure 3.5: Capacitance- and ohmic-corrected cyclic voltammogram data (1.0-1.7 V
versus RHE, 10 mV/s) of a BSCF82 electrode (0.25 mg/cm 2 oxide loading) on the first
cycle (blue line). Also shown are potentiostatic data on the same electrode taken
11
immediately after the first cycle (red X). Measurements taken in O2-saturated 0.1 M
KOH, normalized to geometric electrode area. Figure adapted from published work.1 '..40
Figure 3.6: Selected regions of the data shown in Figure 3.4a. Data are capacitanceand iR-corrected. For each cycle, the average of the forward and backward scan was
mirrored in time to account for current from both forward and backward scans. Figure
adapted from published work." ....................................... .. ... .. .. ... .. .. ... .. .. ... .. .. .. ... .. ... . . . 41
Figure 3.7: Potentiostatic data measured on a single BSCF82 electrode, held at 1.6 V
vs. RHE for 2 h. This electrode did not contain AB carbon. Measurement performed by
Marcel Risch. Figure adapted from published work." ...............................................
42
Figure 3.8: HRTEM and FFTs from approximately 35 x 35 nm areas of (a) asprepared BSCF82, (b) cycled BSCF82, (c) as-prepared LMO, (d) cycled LMO, (e) asprepared LCO, (f) cycled LCO, (g) as-prepared LSC46 and (h) cycled LSC46. Cycling
consisted of five cycles between 1.1-1.7 V vs. RHE in 02-saturated 0.1 M KOH at 10
mV/s. HRTEM performed by Christopher E. Carlton. Figure adapted from published
w o rk ." ............................................................................................................................
43
Figure 3.9: HRTEM and FFTs of (a) as-prepared BSCF82 powder, (b) BSCF82 cycled
5 times, (c) BSCF82 cycled 185 times, (d) as-prepared BSCF46, (e) BSCF46 cycled 5
times, and (f) LSC46 cycled 185 times. Cycles were 1.1-1.7 V vs. RHE at 10 mV/s, in
0-saturated 0.1 M KOH. HRTEM performed by Christopher E. Carlton. Figure
adapted from published w ork." ..................................................................................
44
Figure 3.10: HRTEM and FFT of BSCF82 electrode that was cycled 5 times (a) with
AB carbon and (b) without AB carbon. The amorphous layers are similar and show
that AB carbon does not significantly contribute to the amorphization process. HRTEM
performed by Christopher E. Carlton. Figure adapted from published work."...........44
Figure 3.11: HRTEM and corresponding FFTs of particle surfaces of (a) as-prepared
BSCF82, (b) BSCF82 cycled 5 times, (c) BSCF82 cycled 100 times, (d) BSCF82 held
potentiostatically at 1.7 V vs. RHE for 2 h, and (e) SCF82 cycled 5 times. Cycles were
1.1-1.7 V vs. RHE at 10 mV/s, in 0-saturated 0.1 M KOH. HRTEM performed by
Christopher E. Carlton. Figure adapted from published work.".................................45
Figure 3.12: High angular annular dark field imaging and scanning TEM EDS of
BSCF82 cycled 5 times between 1.1-1.7 V vs. RHE at 10 mV/s, in O-saturated 0.1 M
KOH. The quantitative EDS results are shown in both the images (a) and (b), and
plotted vs. the distance from the particle edge in (c) and (d). Both (a) and (b) are the
same image, with EDS taken at different locations on the particle. A-site atom
concentration (Ba and Sr) is compared to B-site concentration (Co and Fe) for visual
clarity. HRTEM/EDS performed by Christopher E. Carlton. Figure adapted from
p ub lish ed w o rk ." ............................................................................................................
46
12
Figure 3.13: Co 3O4 peaks have been observed in the Raman spectra of both BSCF82
and LSC46, both cycled 5 times between 1.1-1.7 V vs. RHE at 10 mV/s. The multipliers
in the image refer to intensity scaling for visual clarity. Approximately 15% of spectra
exhibited the Co3 O4 -like peaks. These peaks were never observed for as-prepared
samples, nor in LCO samples before or after cycling. Raman spectroscopy performed by
Kelsey A. Stoerzinger. Figure adapted from published work............................47
Figure 3.14: Raman spectra of (a) BSCF82 and (b) LCO before and after cycling 5
and 50 times, (c) BSCF82 after cycling 50 times in the indicated voltage windows, (d)
LMO, (e) LSC46 and (f) BSCF46 before and after cycling 5 and 50 times. Cycling took
place in O2 -saturated 0.1 M KOH at 10 mV/s. Raman spectroscopy performed by
Kelsey A. Stoerzinger. Figure adapted from published work."..........................49
Figure 3.15: Raman spectra (left) of BSCF82 after potentionstatically holding at the
labeled potential until 27.4 mC of charge had passed. Electrochemical data is shown on
the right panel. Times required were 0.06, 1.78, and 24 h for 1.600 (red), 1.475 (green)
and 1.45 (blue) V vs. RHE, respectively. For potentials with appreciable OER current,
the current increased over time, accompanying more drastic changes in the Raman
spectra. Raman spectroscopy performed by Kelsey A. Stoerzinger. Figure adapted from
50
p u b lished w ork ." ............................................................................................................
Figure 3.16: Oxygen evolution current at 1.58 V vs. RHE vs.
passed during a single cycle (1.1-1.7 V vs. RHE, 10 mV/s) in
numbers are labeled. Inset shows a single cycle (5 th) with the
proportional to the cathodic charge. Figure adapted from published
the cathodic charge
0.1 M KOH. Cycle
shaded region being
work." ................. 52
Figure 3.17: Tafel plots for BSCF82 (solid squares, 2 nd cycle; open squares, 1 8 5 th cycle)
and an electrochemically-deposited amorphous Co oxide film (green triangle, 2 "d cycle)
in 0.1 M KOH. Tafel plots constructed from averaged forward-backward scans of cyclic
voltammetry (10 mV/s) and normalized to the total number of Co in bulk or on the
surface. BSCF82 data was gathered on a single electrode while the Co film data was the
average of 5 independent measurements (250 nmol/cm 2 isk loading, 200 nm thick). The
number of cobalt atoms in BSCF82 was estimated from the formula unit, mass loading
and SEM-estimated surface area. For the Co-film, the number of Co ions was estimated
from the Faradaic charge passed during electrodeposition at 0.806 V vs. SCE in 0.1 M
potassium phosphate (pH 7), assuming all charge was due to Co2 to Col oxidation. Co
53
film data was gathered by Marcel Risch. Figure adapted from published work." .....
Figure 3.18: (a) Calculated 0 p-band centers of perovskites BSCF46, BSCF82, SCF82,
LSC46, LCO and LMO, and (b) the overpotentials at 0.25 mA/cm 2 0 x. Stoichiometry in
the calculations is not exact due to the coarse composition grid in the simulated
supercells. Oxides represented with red bars exhibit rapid surface amorphization during
water oxidation, while those represented with blue bars remain highly crystalline. DFT
performed by Yueh-Lin Lee. Figure adapted from published work."..........................55
13
14
List of Tables
Table 3.1: Space groups and lattice parameters of the oxides studied in this work......36
Table 3.2: SEM-estimated surface areas of oxides in this work................................36
15
16
1 Introduction
1.1
Motivation
The risks of anthropogenic climate change have led to intensive research towards
alternative sources of energy to conventional fossil fuels. Greenhouse gas emissions have
steadily increased and are projected to continue to increase in the near future," further
highlighting the need for scientific solutions to the global energy crisis. Clean, renewable
energy sources such as wind and solar energy have attracted significant attention, but
have the disadvantage of being intermittent in nature. One possible avenue to
improving the reliability and practicality of these technologies is a robust and efficient
way of storing intermittent energy as chemical fuel which can be used at a later time. In
particular, molecular hydrogen generated by the electrochemical splitting of water has
received a large amount of focus in the literature.3 -
However, the efficiency of this
process is limited by the sluggish electrochemical kinetics of the oxygen evolution
reaction (OER), which is one half of the overall water splitting reaction (the other half
being the hydrogen evolution reaction, HER).
Commercial grid electrolyzers are attractive for their simplicity and commercial
readiness compared to the direct solar-to-fuel scheme." While larger-scale systems will
start to accumulate ohmic losses that make up a large part of the efficiency losses,
kinetic overpotential loss from the oxygen evolution reaction still remains significant 13
and necessitates the use of precious metal oxides (IrO2, RuO2 ) for acidic electrolytes"
and nickel/nickel oxides which are susceptible to deactivation over time in alkaline
electrolytes."
17
There has been growing interest in the area of solar fuels,"' where sunlight is used
to drive the synthesis of chemical fuel, in a sort of "artificial photosynthesis". For these
applications, where current density is limited by the solar photon flux, highly efficient
catalysts may be particularly important as a way to reduce the photovoltage needed to
sustain a given current, relaxing constraints on semiconductor device performance. The
state-of-the-art OER catalysts typically require around 400 mV of overpotential5'10 to
sustain the current densities achieved in record photoelectrochemical cells.1 5 This is a
non-trivial amount of additional photovoltage for a device to provide.
It is clear that earth-abundant, inexpensive and highly efficient catalysts for the
OER will have an enormous impact on the viability of generating hydrogen from water,
and possibly for the generation of other hydrocarbon fuels as well, where the OER is a
half-reaction. However, the exact mechanism remains relatively poorly understood (i.e.
when compared to HER), even though there have been recent advances in achieving
high-performance catalysts with some first steps towards rational design. 37
01 1
There
remains much room for furthering our fundamental understanding of the reaction
mechanism and the requirements for a highly active, stable catalyst.
1.2 The Oxygen Evolution Reaction: Historical Overview
The OER is a 4 step, 4 electron transfer half-reaction, with the net reaction
typically written in acid as (2H20
2H20 +
02
02
+ 4H+ + 4e-)1
17
and in alkaline as (40H- -
+ 4e ).1'1" The slow kinetics of this reaction have been known-albeit not
fully understood-for nearly a century, along with the fact that the metal electrodes
18
studied frequently in early reports form oxide layers under OER conditions.2 0-"
Hoar
proposed" that the electrochemical irreversibility is due to the permeability of the oxide
film to electrolyte and the resulting self-polarization (not taking into consideration
electrode kinetics), and mentions that a hypothetical catalytically-active stable metal
surface could provide reversible electrochemical behavior-although this is not possible
to test experimentally and is not in agreement with the modern picture of OER
irreversibility based on density functional theory."2 Hoar was also the first to make the
connection that the intersection of the Tafel lines for oxygen evolution and reduction
occurs at 1.23 V vs. reversible hydrogen electrode (RHE), which is the theoretical
thermodynamic potential.2 1 One study by Bockris23 treated the OER (and multi-step
electron transfer reactions in general) in an analogous manner to the much more widely
studied hydrogen electrode, which by that time was starting to be understood to be
governed by the binding strength of intermediate hydrogen species to the catalyst
surface.2" Bockris performed a theoretical analysis of the expected Tafel behavior of a
series of hypothetical reaction pathways, by first expressing rate equations for each step
in the OER, and making several assumptions on the coverage isotherm, species
concentrations and rate-limiting step.2 ' This allowed a table of theoretical Tafel slopes
to be generated for a series of proposed pathways, for comparison to experimental data.
Unfortunately, the Tafel slopes were not unique to each proposed mechanism and
therefore we cannot determine, from this type of analysis, the true mechanism of the
OER-which to this day is still not conclusively understood.
Other early work focused on the attainability of a "reversible" oxygen electrode,
especially in acidic solution with platinum working electrodes and careful elimination of
19
system impurities.18 ,28 -34 Here, "reversible" refers to observing an open-circuit potential
equal to the thermodynamic potential of 1.23 V vs. RHE. This is indeed attainable on
clean platinum electrodes with electrolytes that have been pre-electrolyzed to achieve
low impurity concentrations, however, once appreciable OER current begins, the anodic
oxidation of platinum stops the reversible behavior. 18
35
In general, the open circuit
potential of approximately 1.0 V vs. RHE typically seen on platinum electrodes, is
attributed to a competing process, or mixed potential, between oxygen reduction (for
which Pt is a good catalyst) and platinum oxidation, owing to the extremely slow OER
kinetics.3 4 With this in mind, it is important to distinguish between a surface that can
provide an open circuit potential near the thermodynamically reversible potential and a
surface that can sustain high OER current with low overpotential.
therefore only of interest
Platinum was
as a model surface for fundamental studies of the OER
mechanism in acidic electrolyte.
Other electrodes, such as precious-metals iridium and ruthenium, have also been
investigated. 36
39
Shortly thereafter, the highly-active rutile oxide catalysts IrO (used in
proton exchange membrane electrolyzers) and RuO 2 were discovered.'
4
These materials
have still been subjects of active research in recent times.5 37 42 -4 These oxides reigned as
the most-active OER catalysts until recently,' 0 when transition metal oxides were
demonstrated to be capable of exceeding precious-metal oxides in activity. Transition
metal electrodes
(their oxidized surfaces)
have also been measured, 19'
47
-5
including
perovskite oxides, which will be discussed further in the next section. Typically these
systems are more complicated than the metal surfaces frequently studied in hydrogen
20
evolution/oxidation
can be said
and less
of the exact
nature of the oxygen
electrocatalysis mechanisms on these surfaces based on the electrochemistry data alone."
The mechanism may very well change between different materials. Investigating
transition metal oxide catalysts for OER remains an active area of study.
Recently, focus has increased on amorphous or nano-structured cobalt and nickel
oxides
as
solution. 7' 6 '
OER
7
61
catalysts,
deposited
anodically
in phosphate
or
borate
buffer
Notably, these catalysts have a self-repairing ability in the above-
mentioned buffer solutions and have high catalytic activity even in neutral pH.
Furthermore, when comparing films of different thickness, the activity increases,
suggesting that the oxide films are permeable to electrolyte and that cobalt sites within
the bulk of the material are capable of acting as active sites for the OER. 2 Preliminary
isotope labeling experiments via
and
302
18 0-rich
oxide films revealed significant amounts of
"02
from in-line mass spectrometry, hinting that bulk oxygen may play a role in
the OER mechanism.6 2 Structural studies of amorphous cobalt films via extended X-ray
absorption fine structure (EXAFS) have revealed that the bulk film is composed of
structural motifs or "clusters", consisting of interconnected Co'--oxo cubanes, with
cobalt in edge-sharing octahedra.3 '6" It has been suggested from this structure that the
space between clusters could be filled with cations, anions or water, giving structural
evidence for the possibility of bulk cobalt atoms being active for OER." Furthermore,
the activity seems to scale with the size of these clusters.5
76 5
These amorphous,
electrochemically-deposited oxide films are of importance not only as potentially viable
catalysts in commercial applications, but also for the new fundamental insights that
may be gained through their study.
21
This section ends with a discussion of recent developments in the understanding
of OER kinetics. Perhaps one of the most profound advances in the study of oxygen
electrocatalysis-or electrocatalysis in general-is the "volcano" plot, first introduced
from theoretical considerations by Balandin."
This is a visual representation
of
Sabatier's principle: the optimal catalyst will bind the adsorbed species neither too
strongly nor too weakly." Trasatti presented an experimental volcano plot from the
literature available at the time for the oxygen evolution reaction, with the experimental
OER overpotential at a given current plotted versus the heat of formation from a lower
oxide to a higher oxide (where the metal becomes more oxidized) for the various metals,
effectively giving a measure of the metal-oxygen bond strength." This has led to a
thermodynamic picture of catalysis, where catalytic activity is described in terms of the
binding strength of intermediate species in the overall reaction." In this framework, the
change in Gibbs free energy of adsorbed intermediates must be energetically downhill for
each reaction step in order for the overall reaction to proceed appreciably." In the ideal
case, each intermediate step would result in a change of Gibbs free energy of 1.23 eV,
such that when the electrode potential is 1.23 V vs. RHE, the energy landscape is flat
(Figure 1.1). Some development of this concept was done via density functional theory
calculations of the binding strength of oxygen intermediates on various surfaces.
Rossmeisl et al. found scaling relationships between related oxygen intermediate
species."," This means that the adsorption energy of -00H and -OH, for example, are
not independent. In the context of the thermodynamic framework described above, this
implies that there is a limit to the activity of an OER catalyst." In real systems the
22
changes in Gibbs free energy for different intermediates are not the same, and since they
are not independent of each other they cannot be individually tailored to give the ideal
binding strength (Figure 1.1). It is important to note that this framework does not take
into account transition states, only intermediates bound to the surface, and hence it
does not truly consider kinetics. It remains to be seen whether such scaling relationships
can be overcome by creative catalyst design, such as three-dimensional structures or
multiple types of active site.
V
4.92E
3.69
r
1.23
0
Eapp =
1.23 V
Reaction Coordinate
Figure 1.1: Free energies of OER intermediates on an oxide surface, for the case of zero
applied potential and 1.23 V applied potential. The specific intermediates assumed are not
considered here. The "ideal" catalyst is shown in red, while a more realistic catalyst is shown in
blue. The ideal catalyst has a flat energy landscape at the thermodynamic equilibrium potential
of 1.23 V vs. RHE. Figure adapted from work by Rossmeisl et al.2 ' and Dau et al."
1.3 Perovskite Oxides as Oxygen Evolution Catalysts
Perovskites are a diverse set of compounds typically with formula ABO3 , where A
is typically an alkaline earth or lanthanide metal but can be monovalent, bivalent or
trivalent, and B a transition metal with the ideal cubic structure shown in Figure 1.2.70
The A-site is 12-fold coordinated with oxygen and the B-site is in an octahedral 6-fold
23
coordination with oxygen. Ideally, the B-O distance is a/2, where a is the cubic unit cell
length, and the A-O distance is a/12, giving a Goldschmidt tolerance factor 7 of 1:
t
t=
(rA + ro)
(1.1)
V2(rB + ro)
For perovskites with t ~ 1, and at high temperatures, the structure is usually
cubic. Deviations from t = 1 can result in different distortions of the structure, such as
orthorhombic or rhombohedral. Perovskites are incredibly useful as a set of model
compounds owing to the large variety in chemistry and structure available by changing
the A- and B-site cations and by changing the stoichiometry via alloying. They have
0 72
therefore been used in many studies for oxygen electrocatalysis.1'
76
Typically in electrocatalysis the ion of interest is the B-site, and is usually
assumed to be the active site for oxygen evolution. ''"
From a crystal field theory
perspective, the octahedrally-coordinated B-site's d-state manifold will be split into
several levels, as shown in Figure 1.3. The states of interest for catalysis will be the
antibonding eg and t2g states, since they are typically the occupied states with highest
energy and their filling will roughly determine the strength of the B-O bond.'
77
Figure 1.2: Cubic perovskite structure. The central B-site transition metal (dark blue) is
octahedrally coordinated by oxygen (red), with the A-site alkaline earth or lanthanide metal
(yellow) at the corners in a 12-fold coordination by oxygen.
24
t(,a*, Tc*)
(a*)
ra,
4p
4s
eg
3d
-
(a*)
*
t1g1+t2,
%ti,~(aY,iT)
Energy
e (a)
% tag
(at)
-
,
Figure 1.3: Qualitative molecular orbital diagram for octahedrally-coordinated transition metal
ion. Asterisks denote antibonding states. In bold are the eg and
Figure adapted from Cox."
t2g
states of interest in catalysis.
Numerous perovskite oxides were studied in the early 1980s by Bockris and
coworkers,7 2 7, 3 as well' as Matsumoto et al.79
83
Several possible mechanistic pathways 3
were proposed, attempts were made to correlate the catalytic activity with M-OH bond
strength and number of d-electrons, and the possible importance of different orbital
occupancies was mentioned.7 2 Nickel and cobalt perovskites were typically regarded as
the most active perovskites for oxygen evolution.75 '79 More recently, a scheme was
proposed based on a simplified molecular orbital model that uses the eg orbital
occupancy as a descriptor for perovskite OER activity," similarly to the case for the
oxygen reduction reaction,76 which was itself inspired by the d-band center theory for
metal surfaces.8 4'85 The OER activities of a series of perovskites were found to form a
volcano plot (Figure 1.4) when plotted versus the eg orbital filling, which was
determined via X-ray absorption near edge structure (XANES) and spin states inferred
25
from other studies in the literature. Perovskites with an eg orbital occupancy of
approximately 1 were found to be the most active (e.g. LaNiO3 , LaCoO3 ) and based on
this principle a highly active perovskite catalyst, BaO 5SrO. 5 CoO.8 FeO.20 3 . was discovered,
showing that these descriptors can be used to rationally choose materials as candidates
for promising OER catalysts. From this work, this thesis has extended the study of
some perovskite catalysts, studying further the catalytic activity and surface structure
changes that occur on BaO.5SrO5 Co0 8 Feo.20 3.6and related compounds.
1.4
1 ,
,
.
.
.
,
,
,
/ a 50 pA cmax
NA
LU
LaCoO3
SLaO5CaO,5MnO3
LaMnuCu. 0 O
1.6
OT
1.5
LaMnO3 +,
'
1,8
Ba 0Sr,.Co ,Feo.20 sO4
LaoCa,.,CoO3
La
CoO
NC;
La Ca FeO
10
3
LaMnOa
LaCrO3
-
1(
0.0
0.5
1.0
1.5
2.0
2.5
e 9 electron
Figure 1.4: Volcano plot showing the relationship between OER catalytic activity, defined as
the overpotential at 50 iA cm-2 (estimated true oxide surface area), and the eg orbital occupancy
of the B-site transition metal ion. Figure from previous work. 0
1.4 Scope of this Thesis
This thesis presents an investigation of the surfaces of perovskite oxide catalysts
before and after electrochemical measurements in alkaline electrolyte. The highly active
catalyst BSCF82 was observed to have unusual pseudocapacitance relative to other
perovskites studied.10 Further investigation of the electrochemical behavior of this
26
catalyst and other perovskites is presented. Then, high resolution transmission electron
microscopy
(HRTEM)
shows
that BSCF82
and the closely related
compounds
SrCo 8 Fe, 2 03_ (SCF82) and BaO5 SrO5 Co0 4 Fe. 0 O3 -1 (BSCF46) undergo a structural change
(distinct from corrosion) at their surfaces to an amorphous oxide. This surface change is
concomitant with an increase in the pseudocapacitive current and the OER activity.
The other perovskite catalysts were investigated that remained crystalline under
identical conditions. Raman spectroscopy is shown to be an effective tool for quickly
determining whether an oxide rapidly becomes amorphous at the surface during OER.
The effects of different electrochemical conditions, such as cycling electrodes or holding
potentiostatically, and changing the potential window, are also discussed. Parallels
between the amorphous surface layers and electrochemically-deposited cobalt oxide film
catalysts are drawn. Based on density functional theory calculations, it is shown that
the calculated oxygen p-band center position relative to the Fermi level is a descriptor
for both experimental OER activity and stability in the perovskites studied. Finally,
further interpretations and hypotheses based on the data are discussed along with future
research directions. This work illustrates the importance of understanding the exact
nature of the oxide surface where the OER takes place, in order to further understand
and develop active and stable oxide catalysts for the OER.
27
28
2 Experimental
2.1 Oxide Synthesis and Characterization
Several techniques were utilized in this study for synthesizing perovskite oxides,
including LaCoO 3 (LCO), LaMnO 3 (LMO), La 0,Sr 0 6 CoO3 - (LSC46), SrCo0 8 Fe0 .O 3 (SCF82), and Ba 0 .5 Sr 0 .5 Co1 ,Fe,03_6 (y
0.2, 0.6; BSCF82 and BSCF 46, respectively).
For LCO and LMO a co-precipitation method was used, where metal nitrates (99.98%,
Alfa Aesar) were mixed in the desired molar ratio in deionized water (18 MQ-cm,
Millipore) such that metal concentrations were approximately 0.1 M. This solution was
titrated with 1.2 M tetramethylammonium hydroxide (100%, Alfa Aesar) until a
precipitate formed. This precipitate was then filtered and dried before annealing at
10000 C. LCO was annealed in dry air, while LMO was annealed in Ar.
BSCF82 and BSCF46 were synthesized by a nitrate combustion method.1 " The
alkaline earth and transition metal nitrates (99.9998+%, Sigma-Aldrich) were added in
the desired molar ratio to deionized water at approximately 0.2 M concentration. The
mixture was then heated until most of the water had evaporated, and continued until
sparks were observed in the remaining powder, indicating combustion of the nitrates.
After the sparking had ceased, the powder was annealed in air at 11000 C for 24 h.
SCF82 was synthesized by solid-state reaction. SrCO 3 , Co3 04, and Fe2 0 3 were
mixed in the correct molar ratio, ground with a mortar and pestle and fired in air at
1000' C for 12 h. The powder was then ground again and annealed in air at 1100 C for
24 h with intermediate grindings.
29
LSC46 was synthesized by the Pechini method.8" Similarly to the above methods,
metal nitrates were measured out in the desired molar ratio and added to deionized
water. Citric acid (99%, Sigma-Aldrich) was added in a 5:1 molar ratio of citric acid to
total metal cations. The mixture was stirred at 750 C for 30 min, before adding ethylene
glycol in a 1:1 molar ratio to citric acid. This mixture was covered and stirred at 900 C
for 1 h.
Phase purity was verified via X-ray diffraction with a Rigaku high-power rotating
anode X-ray powder diffractometer with a typical range of 300 to 95' 2,0 under Cr KU
radiation. All oxides used in this study showed no sign of impurity phases.
Particle size analysis was performed using scanning electron microscopy (SEM,
JEOL 6320 FV). Specific area (cm 2/mg) was calculated using a spherical geometry
approximation:8 7
A t 7r 2
( )prd3
6
d2
6
p
d3
pdra
(2.1)
Each sum was calculated over all the individual particles counted in the
micrographs. Adobe Photoshop CS5 was used to select particles using the lasso tool, and
the image analysis tools were used to obtain two-dimensional areas for each selected
particle, from which an effective spherical diameter was obtained for use in equation 2.1.
Powder specific areas determined via this method have been found to agree fairly well
with those obtained by Brunauer, Emmet and Teller (BET) analysis, typically within a
factor of 2-3.1o'88
30
2.2 Electrode Preparation
Electrodes were prepared by a catalyst ink drop-casting technique." Catalyst ink
was prepared by weighing out oxide power and acetylene black (AB) carbon in a 5:1
weight ratio and dispersing in tetrahydrofuran (THF, 99%, Sigma-Aldrich) such that 10
pL of solution contained 0.05 mg of oxide. This dispersion was bath-sonicated for 1 h,
and tip-sonicated for 10 min before adding KOH-neutralized Nafion@ polymer (DE520,
Ion Power, DE) as binder in a 1:1 weight ratio to AB carbon and bath sonicating again
for 2 min. The Nafion solution was prepared by dropping in 0.1 M KOH to the aspurchased Nafion liquid, with a 2:1 volume ratio of Nafion to KOH. Then, 10 gL of ink
was deposited on a glassy carbon rotating disk electrode with a micro-pipette such that
oxide loading was 0.25 mg/cm
2
and left to dry. Electrodes that were prepared for
Raman spectroscopy measurements did not contain AB carbon in order to avoid
background signal.
2.3 Electrochemical Measurements
Electrochemical experiments were conducted using a rotating disk electrode setup
consisting of a borosilicate glass cell, either a saturated calomel reference electrode
(SCE) or Ag/AgCl reference electrode, platinum wire counter electrode and Teflon
electrode holder (all from Pine Instruments) and a Biologic SP-300 potentiostat.
Reference electrodes were calibrated in 0.1 M KOH by bubbling hydrogen gas (ultrahigh purity, Airgas) for 20 minutes and measuring hydrogen oxidation/evolution at a
platinum disk electrode.
The open-circuit potential was then defined as 0 V versus
reversible hydrogen electrode (RHE) after verifying stable, reversible electrochemical
31
behavior. The electrolyte prepared with deionized water (> 18 MQ-cm, Millipore) and
KOH pellets (99.99% purity metals-basis, Sigma-Aldrich) to yield a concentration of 0.1
M. For OER measurements, oxygen gas (ultra-high purity, Airgas) was bubbled for 1520
minutes
measurements,
before
to
electrochemical
ensure
02/H 2 0
measurements,
equilibrium
at
and
1.23
was
V
continued
versus
RHE.
during
Cyclic
voltammetry measurements used a scan rate of 10 mV/s. All measurements used a
rotation rate of 1600 rpm, with the exception of electrodes prepared for Raman
spectroscopy, which used a rotation rate of 900 rpm. There was no observed dependence
of OER current on rotation rate in the potential windows accessed in this study. Highfrequency AC impedance was used to determine the ohmic resistance R of the
electrolyte (typically 45 Q for 0.1 M KOH). Ohmic losses were corrected by subtracting
the ohmic drop iR from the measured potential where i is the measured current and R is
the electrolyte resistance as described above. Potentials corrected for ohmic drop are
denoted as E - iR. OER kinetic currents were capacitance-corrected by taking the
average between positive- and negative-going scans, and the specific activity was
obtained by normalizing the current by the surface area of each oxide as estimated via
SEM measurements described previously.
2.4 Characterization of Electrochemically-Cycled Electrodes
Samples were prepared for transmission electron microscopy (TEM) by removing
the electrode from electrolyte following electrochemical measurements, rinsing briefly
with ethanol, then swabbing a lacy carbon grid over the electrode. Grids were dried and
32
stored in a vacuum oven at 700 C. Samples of the as-synthesized powders were made by
sonicating the powder in THF for 5 minutes and then collecting the powder on a lacy
carbon grid, followed by drying in a vacuum oven at 70* C.
TEM was conducted on a JEOL 2010F instrument operated at 200 keV. The
microscope was equipped with a field-emission electron gun and ultra-high resolution
pole piece, yielding a point-to-point resolution of 1.9
1.
Fourier image analysis was
performed using Gatan Digital Micrograph software v2.01 (Gatan).
Energy-dispersive X-ray spectroscopy (EDS) was performed by operating the TEM
in scanning mode using an electron beam of approximately 0.7 nm. INCA (Oxford
Instruments) software was used for controlling beam position and for data analysis. Live
acquisition time was 60 s for each EDS spectrum, with the beam placed at various
positions on the catalyst surface or bulk, with a spatial resolution of approximately 2
nm owing to sample drift during data acquisition.
Raman spectroscopy performed on electrodes after electrochemical measurements
was done by removing the electrode from electrolyte, drying with compressed air and
using the electrode directly under the microscope of the LabRAM HR800 (Horiba Jobin
Yvon). A 633 nm He/Ne laser at 0.2 or 2 mW was used with a 50x objective (giving
spot size of approximately 30 Iu2 ), at room temperature. Sampling time was adjusted
per sample to yield approximately the same signal-to-noise ratio for each measurement
(6x30 s for LMO, 6x60 s for LCO, 6x60 s for BSCF82 (0.7-1.0 V potential range),
3x120 s for BSCF82 (all other potential ranges), and 6x120 s for LSC46). No change in
the Raman peak intensity was observed over the course of the 3-12 min sampling time.
Spectra were collected at six or more locations on each electrode.
33
2.5 Density Functional Theory Calculations
Density
functional theory
(DFT) calculations used
the Vienna
plane-wave method.89 91
simulation package
(VASP)
Exchange-correlation
was treated with the Perdew-Wang-91
approximation (GGA)
with the Project-Augmented
with soft O
ab initio
generalized gradient
oxygen pseudopotential and the Hubbard U
correction.92,9 Energy convergence was within 3 meV per perovskite formula unit from a
Monkhorst-Pack 4x4x4 k-point mesh. The calculations are performed using the
simplified spherically averaged approach,9 4 where Uff is applied to d electrons ( Uj =
Uoulomb
-
exchange J) with values of 4, 4, and 3.3 eV for Mn,
Fe, and Co,
respectively. 95 ' 96 Typically, 2x2x2 perovskite supercells were used for simulations and
fully stoichiometric compounds were used (no oxygen vacancies). For perovskites with
disordered cations (i.e. BSCF46, BSCF82, SCF82 and LSC46) a special quasirandom
structure (SQS) approach was used.97 Several sampled configurations' calculated 0 pband centers were used in a weighted average (weighted by the Boltzmann factor at 300
K for the energy of the sampled configurations) to achieve an effective 0 p-band center.
34
3 Results and Discussion
3.1 Material Characterization
X-Ray diffraction (XRD) was performed on all oxides, namely LaCoO, (LCO),
LaMnO 3 (LMO), LaOSrO6CoO 3 - (LSC46),
,Fe,0
3-
SrCo0 8FeO.2036 (SCF82), and Ba,0 5Sr 0 5Co1
(y = 0.2, 0.6; BSCF82 and BSCF 46, respectively), to ensure phase purity. The
resulting diffractometry results are shown in Figure 3.1, with relevant extracted lattice
parameters summarized in Table 3.1. All XRD data was gathered from pure oxide
powder with the exception of LSC46, which included Si as a reference (Figure 3.2).
75
AA
60 --6
S4LSC46
-
LMO
LCO
30
BSCF46
.l|15 -
BSCF82
B-A.SCF82
00-
4
3
2
1
d-spacing (A)
Figure 3.1: X-ray diffraction data for the oxides studied in this work. LMO, LCO data were
gathered by Jin Suntivich. LSC46 data were gathered by Ethan J. Crumlin. SCF82 data were
gathered by Alexis Grimaud.
35
Experimental
La0ASr ,,CoO,
I
II
IIIIII
Si
4.0
3.0
2.0
I
1.0
d-spacing (A)
Figure 3.2: Experimental XRD peak locations (red) with database peak positions for LSC46
(gray) a nd Si (green). Si was included to give reference diffraction peaks.
T able 3.1: Space groups and lattice parameters of the oxides studied in this work.
Lattice Parameters (A)
b
c
Material
Space Group
LMO
Pnma
5.66
7.22
5.53
LCO
R3c
5.44
5.44
13.09
LSC46
Pm3m
3.84
3.84
3.84
BSCF46
Pm3m
3.96
3.96
3.96
SCF82
Pm3m
3.86
3.86
3.86
BSCF82
Pm3m
3.99
3.99
3.99
a
C xide surface area estimations were calculated based on scanning electron
microscopy (SEM; see Chapter 2.1). Calculated values of As are tabulated in Table 3.2.
Table 3.2: SEM-estimated surface areas of oxides in this work.
Material
Specific Area (m 2 .g')
LMO
0.6
36
LCO
0.7
LSC46
1.2
BSCF46
0.1
SCF82
0.1
BSCF82
0.2
3.2 Electrochemical Measurements
Cyclic voltammetry was performed on LCO, LMO, LSC46, SCF82, BSCF46 and
BSCF82, with the 1t and
5 th
cycles shown in Figure 3.3. The OER activity, defined as
the overpotential required for a given current density normalized to true oxide surface
area, increases as expected,1 " with LMO < LCO < LSC46
~ BSCF46 < SCF82 ~
BSCF82. The cyclic voltammograms of LCO, LMO, and LSC46 all reached a steadystate after the first five cycles, after which no significant changes in the OER activity
were observed. In contrast, BSCF82 exhibited a drastic change in both pseudocapacitive
current and OER activity. OER activity increased steadily with cycling up to around
the 50 cycle
mark
(Figure 3.4a), 'after
which it reached
a steady-state with
approximately 4 times greater OER current at 1.6 V versus reversible hydrogen
electrode (RHE) than that of the first cycle. In contrast, LSC46 cycled up to 100 times
did not undergo any significant changes in OER behavior (Figure 3.4b).
Potentiostatic measurements were performed to observe any changes versus time,
and also to verify that the cyclic voltammetry measurement with the appropriate data
analysis yielded accurate values for the OER activity. Figure 3.4c shows potentiostatic
measurements at four applied potentials, on separate electrodes from those used for
cyclic voltammetry. In Figure 3,4d, any discrepancy between the potentiostatic (solid
marker) and cyclic voltammetry data (hollow marker) is explained by the experimental
error in the electrode deposition and not by the technique used, as seen in Figure 3.5.
37
(a)
(b)
E
E
Ec
1.2
1.3 1.4 1.5 1.6
E -iR (V vs. RHE)
1.7
1.2
1.3 1.4 1.5 1.6
E -iR (V vs. RHE)
1.7
1.2
1.3 1.4 1.5 1.6
E-iR (Vvs. RHE)
1.7
1.2
1.3 1.4 1.5 1.6
E -iR (V vs. RHE)
1.7
(d)
10
1
0
E
5
E
0
1.2
(e)
1.3 1.4 1.5 1.6
E-iR (Vvs. RHE)
1.7
30
30
20
S20
E
E
E 10
0
4~10
E
0
0
1.2
1.3 1.4 1.5 1.6
E -IR (V vs. RHE)
1.7
Figure 3.3: Cyclic voltammetry for (a) LMO, (b) LCO, (c) LSC46, (d) BSCF46, (e) BSCF82
and (f) SCF82. Number labels n indicate the nth cycle of cyclic voltammetry (1.1-1.7 V versus
RHE, 10 mV/s). All measurements were performed in 0 2-saturated 0.1 M KOH, corrected for
ohmic losses and normalized to the SEM-estimated surface area of oxides, and used an oxide
loading of 0.25 mg/cm2 supported on glassy carbon. Figure adapted from published work."
38
40
(a)
(b)
30
10
20
c~a~
E
10
0
5
0
-10
0
-20
1.2
1.2 1.3 1.4 1.5 1.6 1.7
E - iR (V vs. RHE)
(C)
1.3 1.4 1.5 1.6 1.7
E - iR (V vs. RHE)
(d) 1.75
100
jO*1.70-
10
W
1
0.1
>q:
1.65-
1E-4
80
Time (min)
1
1
Is'
515
0
1.601
' 1.55-
0.01
1
LCO
1.50.1
1E-3 0.01
5
SCF82
1
0.1
I (mA cm~O)
50
A
100
BSCF82
10
100
Figure 3.4: Heavy cycling of (a) BSCF82 and (b) LSC46; (c) Potentiostatic measurements on
BSCF82, LSC46, LCO and LMO; (d) Tafel plot containing cyclic Voltammetry at the th cycles
(hollow markers) and potentiostatic data (solid markers) for comparison. Figure adapted from
published work. 1
I
39
15
X
Potentiostatic
Capactance-corrected CV
-0
E
0
.
1.0
I
I
.
a
I
1.2
1.4
1.6
E - iR (V vs. RHE)
Figure 3.5: Capacitance- and ohmic-corrected cyclic voltammogram data (1.0-1.7 V versus
RHE, 10 mV/s) of a BSCF82 electrode (0.25 mg/cm
2
oxide loading) on the first cycle (blue
line). Also shown are potentiostatic data on the same electrode taken immediately after the first
cycle (red X). Measurements taken in 0 2-saturated 0.1 M KOH, normalized to geometric
electrode area. Figure adapted from published work."
In verifying that the current observed is attributable to OER, we first note that
it has previously been observed from rotating ring-disk experiments that current in this
potential range is indeed almost 100% OER."
To rule out corrosion or oxidation of
cobalt/iron in BSCF82 as being significant contributions to the observed current, the
ratio of electrons passed during the observed oxidation current to the total number of
cobalt and iron atoms present in BSCF82 was calculated. To obtain the total charge
passed that is attributable to OER, the capacitance- and iR-corrected current was
integrated (using the data from Figure 3.4a). To include current from both forward and
backward scans, the average of the forward and backward scans was mirrored in time
40
for each cycle, as seen in Figure 3.6. Using this method, the ratio of electrons passed to
total
cobalt
and
iron
is over
atoms
potentiostatic/chronoamperometric
100
for
both
cyclic
voltammetry
and
(Figure 3.7) measurements. The ratio of oxygen
evolved to total oxygen present in the BSCF82, also calculated from amount of Faradaic
charge transfer, is approximately 10 for both cyclic voltammetry and potentiostatic
measurements. Because the catalytic activity was not observed to decrease during the
course of any experiments, dissolution or corrosion current is negligible when compared
with the OER current.
12
CM8
E
o4
E
Capacitance conected and
0
300
7500
7800
t(s)
Figure 3.6: Selected regions of the data shown in Figure 3.4a. Data are capacitance- and iRcorrected. For each cycle, the average of the forward and backward scan was mirrored in time to
account for current from both forward and backward scans. Figure adapted from published
work."
41
1.8
.
.
1.6
--
1.4
E
1.2
E
1.0
0.8
0.6
0
---2000 4000 6000
Time (s)
Figure 3.7: Potentiostatic data measured on a single BSCF82 electrode, held at 1.6 V vs. RHE
for 2 h. This electrode did not contain AB carbon. Measurement performed by Marcel Risch.
Figure adapted from published work."
3.3 Transmission Electron Microscopy
High-resolution transmission electron microscopy (HRTEM) was used to examine
the surfaces both as-prepared powder and electrodes that had been cycled five times
from 1.1 to 1.7 V versus RHE. The surfaces of the as-prepared oxide powders were
highly crystalline; the surfaces of BSCF82 had very few thin amorphous regions. The
fast Fourier transforms (FFTs) of all the HRTEM images of the as-prepared particle
surfaces can all be indexed to the bulk crystal structure of the corresponding oxide
(Figure 3.8). After cycling, amorphous regions around 10 nm thick were present on the
surfaces of BSCF82, SCF82 and BSCF46. For BSCF82 cycled 185 times, the particles
became amorphous to the extent that could be observed with HRTEM (Figure 3.9c).
42
HRTEM was performed on electrodes with and without AB carbon, with no significant
differences between the two (Figure 3.10). In contrast, the surfaces of LMO, LCO and
LSC46 remained crystalline after cycling, comparable to the as-prepared samples (Figure
3.8c-h). This does preclude instability in longer time scales or different voltage regimes.
Figure 3.8: HRTEM and FFTs from approximately 35 x 35 nm areas of (a) as-prepared
BSCF82, (b) cycled BSCF82, (c) as-prepared LMO, (d) cycled LMO, (e) as-prepared LCO, (f)
cycled LCO, (g) as-prepared LSC46 and (h) cycled LSC46. Cycling consisted of five cycles
between 1.1-1.7 V vs. RHE in 0 2-saturated 0.1 M KOH at 10 mV/s. HRTEM performed by
Christopher E. Carlton. Figure adapted from published work."
43
Figure 3.9: HRTEM and FFTs of (a) as-prepared BSCF82 powder, (b) BSCF82 cycled 5
times, (c) BSCF82 cycled 185 times, (d) as-prepared BSCF46, (e) BSCF46 cycled 5 times, and
(f) LSC46 cycled 185 times. Cycles were 1.1-1.7 V vs. RHE at 10 mV/s, in 0 2-saturated 0.1 M
KOH. HRTEM performed by Christopher E. Carlton. Figure adapted from published work."
Figure 3.10: HRTEM and FFT of BSCF82 electrode that was cycled 5 times (a) with AB
carbon and (b) without AB carbon. The amorphous layers are similar and show that AB carbon
does not significantly contribute to the amorphization process. HRTEM performed by
Christopher E. Carlton. Figure adapted from published work."
44
Analysis of the FFT patterns from HRTEM images collected from the surface
amorphous regions of BSCF82 revealed a bright diffuse ring that corresponds to a
spacing of approximately 2.8
A in real
space (Figure 3.11b-c). This spacing, which is not
present in the interatomic distances of cubic perovskite (containing only corner-shared
transition
metal
octahedra),
matches
the
cobalt/iron
spacing
in
edge-sharing
octahedra.98 '99 Energy-dispersive X-ray spectroscopy (EDS) performed in the surface
amorphous regions of BSCF82 revealed a decrease in the concentrations of the A-site
cations Ba 2 ' and Sr2+, which suggests A-site cation leaching during cycling (Figure
3.12).
Ba/Sr (A-site)
Co/Fe (B-sie)
Figure 3.11: HRTEM and corresponding FFTs of particle surfaces of (a) as-prepared BSCF82,
(b) BSCF82 cycled 5 times, (c) BSCF82 cycled 100 times, (d) BSCF82 held potentiostatically at
1.7 V vs. RHE for 2 h, and (e) SCF82 cycled 5 times. Cycles were 1.1-1.7 V vs. RHE at 10
mV/s, in Or-saturated 0.1 M KOH. HRTEM performed by Christopher E. Carlton. Figure
adapted from published work."
45
c)
100
d)
100
0 -
80
;6040
60
40
20
20
10
0
Distance from Surface (nm)
Distance from Surface (nm)
Figure 3.12: High angular annular dark field imaging and scanning TEM EDS of BSCF82
cycled 5 times between 1.1-1.7 V vs. RHE at 10 mV/s, in O 2-saturated 0.1 M KOH. The
quantitative EDS results are shown in both the images (a) and (b), and plotted vs. the distance
from the particle edge in (c) and (d). Both (a) and (b) are the same image, with EDS taken at
different locations on the particle. A-site atom concentration (Ba and Sr) is compared to B-site
concentration (Co and Fe) for visual clarity. HRTEM/EDS performed by Christopher E.
Carlton. Figure adapted from published work."
Amorphous
regions were also found in BSCF82
electrodes that were held
potentiostatically at 1.7 V versus RHE for 2 h (Figure 3.11d). Although the total
amount of charge passed was much higher during this experiment than in the case of
cycling 5 times, the thickness of these regions was comparable to the 5x cycled case.
Interestingly, this suggests that cycling the potential dramatically increases the rate of
46
the amorphization process. Also, the in the FFT of the HRTEM images of the
potentiostatically held electrode, a diffuse ring with real space spacing of approximately
3.3
A
was observed, which corresponds to interatomic distance of cations in spinel
Co3 O4 . This could indicate the formation of spinel-like motifs containing both octahedral
and tetrahedral environments for cobalt and iron cations, as reported for BSCF82 at
intermediate temperatures.
The existence of spinel-like motifs is further supported by
Raman spectroscopy, where approximately 15% of all spectra contained peaks similar to
those present in spinel Co 3O4 (Figure 3.13).
LSC46 (20X)
BSCF82 (4x)
CO304
1000
800
600
400
1
Raman shift (cm )
Figure 3.13: CoO 4 peaks have been observed in the Raman spectra of both BSCF82 and
LSC46, both cycled 5 times between 1.1-1.7 V vs. RHE at 10 mV/s. The multipliers in the
image refer to intensity scaling for visual clarity. Approximately 15% of spectra exhibited the
200
CoaO 4 -like peaks. These peaks were never observed for as-prepared samples or in LCO samples
before or after cycling. Raman spectroscopy performed by Kelsey A. Stoerzinger. Figure adapted
from published work."
47
3.4 Raman Spectroscopy and Potential-Dependent Amorphization
Raman spectroscopy can be used as a rapid tool for probing the amorphization of
oxide surfaces after it has been established (e.g. via HRTEM) that such a process
occurs, and the oxides in question are Raman active. The Raman spectra of BSCF82
prior to cycling exhibits a broad peak at around 675 cm-1, attributed to internal motion
of oxygen within the cobalt and iron octahedra."' It is of interest to investigate the
conditions under which amorphization occurs, by varying the potential window used.
When BSCF82 was cycled between 1.1 and 1.7 V versus RHE as in the electrodes
examined via HRTEM, the vibrational mode of BSCF82 mentioned above broadens
after 5 cycles and was absent after 50 cycles. This is in good agreement with the
changes in the electrochemical data and the amorphization observed via HRTEM. Since
the information depth of Raman spectroscopy is estimated to be larger than 100 nm, the
lack of Raman signal after 50 cycles indicates that the amorphous layer is likely larger
than
100
nm, beyond
the
thicknesses
that
can
be
observed
with
HRTEM
(approximately 20 nm). In contrast, the Raman spectra for LMO, LCO and LSC46
(Figure 3.14) have negligible changes after cycling. Of particular interest are the peaks
at 420 and 600 cm-1 for LCO, which have been attributed to bending and breathing
modes of the oxygen cage of Eg and A2g symmetry, respectively.1
2
These peaks remained
visible after 50 cycles (Figure 3.14b), showing that there is no significant change in the
local symmetry of LCO when cycled under identical conditions to those of BSCF82.
The effect of changing the cyclic voltammetry window was investigated by
cycling BSCF82 50 times (at 10 mV/s) in three voltage ranges: 0.7-1.0 V, 1.1-1.5 V, and
48
1.5-1.7 V versus RHE. No significant changes to the Raman spectra were observed for
the 0.7-1.0 V and 1.1-1.5 V voltage ranges; the 1.5-1.7 V range, showed loss of Raman
scattering accompanying the formation of a thick amorphous layer (Figure 3.14c). This
is potential-dependent
suggests that amorphization
and occurs most rapidly at
potentials above 1.5 V vs. RHE, where OER current is appreciable.
BSCF2
(b)
LCO
(a)
BSCF82
50cS~c1.5-1.7V
5c
5C1.1-1.5V
AP
AP
0.7-1.0 V
300 500 700 900
300 500 700 900
Rmuan shif (oM )
300 500 700 900
(d)
(0)
(4
LMO
A
S.
50c
c
sc
Li
LSC46
50 C
5c
OOO/
C
AP
AP
300 500 700 900
BSCF46
300 500 700 900
Raman shift (am4)
300 500 700 00
Figure 3.14: Raman spectra of (a) BSCF82 and (b) LCO before and after cycling 5 and 50
times, (c) BSCF82 after cycling 50 times in the indicated voltage windows, (d) LMO, (e) LSC46
and (f) BSCF46 before and after cycling 5 and 50 times. Cycling took place in 0 2-saturated 0.1
M KOH at 10 mV/s. Raman spectroscopy performed by Kelsey A. Stoerzinger. Figure adapted
from published work."
Various potentiostatic conditions were also investigated, comparing different
voltages with a fixed amount of charge passed. Similarly to the BSCF82 after cycling,
reduction of Raman scattering was observed at OER potentials (Figure 3.15). Oxidation
49
current was also seen to increase over time; only when potential was held in the OER
region. At 1.45 V versus RHE for 24 h, no strong decrease in Raman signal was
observed, indicating that any amorphization layer is extremely small. OER potential
ranges may be essential for the rapid amorphization observed in BSCF82 and closely
related perovskites. Thus it is possible that OER may kinetically permit or enhance the
amorphization process, although this requires further investigation.
1.0
0.8
1.600 V
0.6
-
0.4
E
-
0.2
1.475 V
0.0
a I
0
1.450 V
a
100
Time (s)
I
-
200
300 500 700 900
Raman shift (cm4)
Figure 3.15: Raman spectra (left) of BSCF82 after potentiostatically holding at the labeled
potential until 27.4 mC of charge had passed. Electrochemical data is shown on the right panel.
Times required were 0.06, 1.78, and 24 h for 1.600 (red), 1.475 (green) and 1.45 (blue) V vs.
RHE, respectively. For potentials with appreciable OER current, the current increased over
time, accompanying more drastic changes in the Raman spectra. Raman spectroscopy performed
by Kelsey A. Stoerzinger. Figure adapted from published work."
50
3.5 Amorphization and Pseudocapacitive Current
The simultaneous increase in pseudocapacitive
current, OER current,
and
amorphous layer thickness leads to the possibility that the amorphous regions may be
accessible to electrolyte.03
105
This would potentially increase the effective number of
active sites for OER, and would also allow redox of the B-site cations to occur in the
bulk of the material, possibly explaining the change in pseudocapacitive current. A
rough
measure of the electrochemically
active surface
area was performed by
determining the cathodic charge passed during the negative-going scan of cyclic
voltammetry, which would be proportional to the pseudocapacitance and the effective
surface area. When plotted versus the OER current at a given overpotential (Figure
3.16), there is a linear trend, where the OER activity increases four to five times and is
correlated with a factor of seven to eight increase in the pseudocapacitive cathodic
charge passed. This lends some support to the claim that amorphization leads to an
increase in the electrochemically active surface area. It should also be noted that on a
per-volume basis the number of cobalt and iron ions present in the amorphous regions
increases faster with region growth than the OER current, and thus not all of the
transition metal cations are electrochemically available as active sites. These findings do
not alter any previous claims of the high catalytic activity of BSCF82," where the
activity was typically extracted from the second or third cycle where the amorphous
layer is still thin. The maximum increase in electrochemically active surface area from
such a layer is less than a factor of two. It has also been recently shown that related
51
cobalt perovskites can sustain OER activities even higher than BSCF82 without
becoming amorphous.06
M: 30
C6
5th cycle
> (2025
>
0 20
0
810.1
100 .
oj
1.0 1.2 1.4 1.6
E-iR(Vvs. RHE)
15
10
''
w
5
C
2
X
0
E
O-'
5
20
40
30
0
-
50
'0
15 20
10
60
80 100 120
Cathodic Charge (mC)
Figure 3.16: Oxygen evolution current at 1.58 V vs. RHE vs. the cathodic charge passed
during a single cycle (1.1-1.7 V vs. RHE, 10 mV/s) in 0.1 M KOH. Cycle numbers are labeled.
Inset shows a single cycle (5 th) with the shaded region being proportional to the cathodic charge.
Figure adapted from published work."
An interesting comparison arises when considering the highly active OER
catalysts formed from electrodeposited cobalt oxide films. When normalizing the OER
current of BSCF82 to the total number of cobalt atoms-taken as the lower limit of the
intrinsic OER activity per cobalt-it was found to be comparable to that reported for
the electrochemically deposited films,1, 59 6, 2,65 as shown in Figure 3.17. Recently it has
been shown that the OER current per cobalt atom decreases with increasing "cluster"
size of an extended cobalt oxide motif, which is a result of reduced utility of cobalt for
OER.
52
7
Another recent report has investigated the local structures of the cobalt cations
via extended X-ray absorption fine structure (EXAFS), and has shown that during
cycling, perovskite oxides BSCF82 and SCF82 both show an increase in the interatomic
spacing associated with edge-shared cobalt octahedra, and begin to show similarity with
the Fourier-transformed EXAFS of the electrochemically deposited amorphous cobalt
oxide films.10 7 Further investigation of the nature of the active sites in the amorphous
regions of these perovskites and their relationship to the electrochemically deposited
oxides is needed to obtain a fuller picture of how these catalysts function.
0.40
A Cofilm,2nd
* BSCF82, 2nd
O BSCF82, 185th
N
0
,;0.35
surface
bulk
i~e~I
U.30
I
IE-18
j
IE-16
IE-14
I (mA/# Co)
Figure 3.17: Tafel plots for BSCF82 (solid squares, 2 "dcycle; open squares, 1 8 5 th cycle) and an
electrochemically-deposited amorphous Co oxide film (green triangle, 2 nd cycle) in 0.1 M KOH.
Tafel plots constructed from averaged forward-backward scans of cyclic voltammetry (10 mV/s)
and normalized to the total number of Co in bulk or on the surface. BSCF82 data was gathered
on a single electrode while the Co film data was the average of 5 independent measurements
(250 nmol/cm 2 di,k loading, 200 nm thick). The number of cobalt atoms in BSCF82 was estimated
from the formula unit, mass loading and SEM-estimated surface area. For the Co-film, the
number of Co ions was estimated from the Faradaic charge passed during electrodeposition at
0.806 V vs. SCE in 0.1 M potassium phosphate (pH 7), assuming all charge was due to Co2 + to
Co"* oxidation. Co film data was gathered by Marcel Risch. Figure adapted from published
work."
53
3.6 0 p-band Center: A Descriptor from Density Functional Theory
Thus far only the amorphization and its effect on catalytic activity have been
discussed. In the following, a potential descriptor for OER activity and amorphization is
reported based on density functional theory calculations. It has recently been reported
that the oxygen p-band center, as calculated from density functional theory, acts as a
descriptor for several physical parameters relating to oxygen anions in the perovskite
structure such as the oxygen vacancy formation energy, oxygen mobility, and oxygen
surface exchange coefficient.9" This is turn leads the p-band center to describe the high
temperature ORR activity of perovskite oxides as cathodes for solid oxide fuel cells.
BSCF82, SCF82, and BSCF46 all have 0 p-band centers much closer to the Fermi level,
allowing oxygen redox to occur more easily, leading to Fermi level pinning of the 0 pband. The calculated 0 p-band centers for the oxides in this work, along with their
OER overpotentials at 0.25 mA/cm'0 x, are shown in Figure 3.18. Oxides with an 0 pband center higher than approximately -2.2 eV relative to the Fermi level become
amorphous during the OER in pH 13 electrolyte, while those with lower 0 p-band
centers do not. Since LSC46 has an 0 p-band center at -2.2 eV, it is likely that an oxide
at the threshold of becoming amorphous during OER has an 0
p-band center
somewhere between that of LSC46 and SCF82. Indeed, there has been further study of
the 0 p-band centers of related double-perovskite compounds that remain stable during
OER with high activity, possessing 0 p-band centers of approximately -1.8 eV relative
to the Fermi level.106
54
Because a high 0 p-band center has been shown to provide higher oxygen lattice
mobility and higher oxygen vacancy concentration in BSCF82,
08
BSCF46 and SCF82,
changes in oxygen stoichiometry at the surface, whether from surface oxygen exchange
or bulk vacancy migration, may play an important role in the amorphization process, or
even in the OER mechanism itself. Further work is needed to understand the intimate
relationship between lattice oxygen, structural stability and catalytic activity in alkaline
solution.
(a)
(b)
BSCF46
BSCF82
SCF82
LSC46
LCO
LMO
-3.0 -2.5 -2.0 -1.5 -1.0 -0.5 0.0
0 p-band center relative to EF (OV)
LLLL
0.0
0.1
0.2
0.3
0.4
0.5
q (V vs. RHE)
Figure 3.18: (a) Calculated 0 p-band centers of perovskites BSCF46, BSCF82, SCF82, LSC46,
LCO and LMO, and (b) the overpotentials at 0.25 mA/cm 20 ,. Stoichiometry in the calculations
is not exact due to the coarse composition grid in the simulated supercells. Oxides represented
with red bars exhibit rapid surface amorphization during water oxidation, while those
represented with blue bars remain highly crystalline. DFT performed by Yueh-Lin Lee. Figure
adapted from published work."
55
56
4 Conclusions and Looking Forward
In this thesis, the amorphization of highly active BSCF82 and related compounds
These oxides rapidly became
BSCF46 and SCF82 during OER was described.
amorphous at the surface when cycled in the OER potential region, with slower
amorphization occurring
for potentiostatically-held
electrodes.
Amorphization was
accompanied by leaching of A-site cations from the surface and considerable increases in
the pseudocapacitive and OER currents. Amorphization also lead to the apparent
A
atomic spacing of 2.8
from the ring observed in the FFTs of HRTEM images of the
particle surfaces, indicating the possible formation of edge-sharing transition metal
octahedra. All three of these oxides had similarly positioned 0 p-band centers as
calculated by density functional theory, being the highest relative to the Fermi level
compared to all the other oxides studied in this work. The perovskites LSC46, LCO and
LMO, which had lower 0 p-band centers relative to the Fermi level, did not become
amorphous and remained crystalline under identical electrochemical conditions.
There are several interesting research directions that arise from the results
presented earlier. First is the parallel between the amorphous layer grown on BSCF82,
BSCF46 and SCF82, and amorphous electrochemically-deposited cobalt films. Further
work has recently been done via EXAFS on BSCF82, SCF82 and the amorphous films,
showing
that
in
the
Fourier-transformed
EXAFS
interatomic
distance
peaks
corresponding to Co-Co distances in edge-sharing octahedra appear and grow with
electrochemical cycling in the OER region."" The possible similarity between the
amorphous cobalt films and the amorphous layers formed on perovskites is what first
57
posed the question as to whether or not the perovskites' amorphous films were
permeable to electrolyte, giving access to bulk cobalt as active sites for OER. It has
already been reported that oxides formed on cobalt metal electrochemically are thought
to be hydrous, and the film growth cyclic voltammetry shows an increase in both
pseudocapacitance and OER current as the film becomes thicker." It had also been
suggested that hydrous oxide films could be an example of three-dimensional catalysis,
where the accessibility of bulk metal cation sites to electrolyte provide a higher
electrochemically active surface area.4 7 109 Similar findings had also been reported for
iridium oxide cycled in acidic solution.n'0
or not the surface oxide
2
Further experiments to determine whether
films contain hydrous species such as hydroxide or
oxyhydroxide and how far into a heavily amorphized film the hydrous species exist, are
needed. Probing the subtle differences between different types of amorphous cobalt (or
other transition metal) oxides, such as studying the effect of any remaining A-site
cations in the amorphous layers, is also of interest to further understand how to
optimize such materials for OER.
Another interesting point is the fact that currently, the most active cobalt-based
oxide OER catalysts were primarily researched initially as high-temperature oxygen
reduction cathode materials for solid-oxide fuel cells (SOFCs) .113,114 It seems clear that
the nature of the transition metal-oxygen bond is of vital importance for OER catalysis.
Differential electrochemical mass spectrometry (DEMS) combined with isotope labeling
of oxygen could be of interest in determining to what degree lattice oxygen participates
in the OER. The degree to which oxygen vacancies are mobile in the bulk at room
58
temperature, and whether oxygen can be exchanged at the oxide-electrolyte interface
could also be of interest to study, as changing oxygen stoichiometry locally near the
surface could have profound effects on oxide stability and OER activity. Investigating
the effects of pH and other electrolytes may also yield new insights. Overall, the many
questions surrounding the interrelationship between lattice oxygen, OER activity, and
oxide stability and its implications for permeability to electrolyte mean that there may
be many nascent advances waiting to be made in the field of OER catalysis on oxides.
59
60
5 References
1.
Birol, F. World Energy Outlook 2011. (Paris, 2011).
2.
Conti, J., Holtberg, P. & Doman, L. E. International Energy Outlook 2011.
(Washington, 2011).
3.
Dinca, M., Surendranath, Y. & Nocera, D. G. Nickel-borate oxygen-evolving
catalyst that functions under benign conditions. Proceedings of the National
Academy of Sciences of the United States of America 107, 10337 (2010).
4.
Gray, H. B. Powering the planet with solar fuel. Nature Chemistry 1, 7 (2009).
5.
Lee, Y., Suntivich, J., May, K. J., Perry, E. E. & Shao-Horn, Y. Synthesis and
Activities of Rutile IrO 2 and RuO 2 Nanoparticles for Oxygen Evolution in Acid
and Alkaline Solutions. The Journal of Physical Chemistry Letters 3, 399 (2012).
6.
Lewis, N. S. & Nocera, D. G. Powering the planet: chemical challenges in solar
energy utilization. Proceedings of the National Academy of Sciences of the United
States of America 103, 15729 (2006).
7.
Kanan, M. W. & Nocera, D. G. In situ formation of an oxygen-evolving catalyst
in neutral water containing phosphate and Co2 . Science 321, 1072 (2008).
8.
Trotochaud, L., Ranney, J. K., Williams, K. N. & Boettcher, S. W. Solution-Cast
Metal Oxide Thin Film Electrocatalysts for Oxygen Evolution. Journal of the
American Chemical Society 134, 17253 (2012).
9.
Smith, R. D. L. et al. Photochemical Route for Accessing Amorphous Metal Oxide
Materials for Water Oxidation Catalysis. Science 340, 60 (2013).
10.
Suntivich, J., May, K. J., Gasteiger, H. A., Goodenough, J. B. & Shao-Horn, Y. A
Perovskite Oxide Optimized for Oxygen Evolution Catalysis from Molecular
Orbital Principles. Science 334, 1383 (2011).
11.
May, K. J. et al. Influence of Oxygen Evolution during Water Oxidation on the
Surface of Perovskite Oxide Catalysts. The Journal of Physical Chemistry Letters
3, 3264 (2012).
12.
Zeng, K. & Zhang, D. Recent progress in alkaline water electrolysis for hydrogen
production and applications. Progress in Energy and Combustion Science 36, 307
(2010).
13.
Choi, P. A simple model for solid polymer electrolyte (SPE) water electrolysis.
Solid State Ionics 175, 535(2004).
14.
Rasten, E., Hagen, G. & Tunold, R. Electrocatalysis in water electrolysis with
solid polymer electrolyte. Electrochimica Acta 48, 3945 (2003).
61
15.
Khaselev, 0. & Turner, J. J. A. A Monolithic Photovoltaic-Photoelectrochemical
Device for Hydrogen Production via Water Splitting. Science 280, 425 (1998).
16.
Dau, H. et al. The Mechanism of Water Oxidation: From Electrolysis via
Homogeneous to Biological Catalysis. ChemCatChem 2, 724 (2010).
17.
Conway, B. & Salomon, M. Electrochemical reaction orders: Applications to the
hydrogen- and oxygen-evolution reactions. Electrochimica Acta 9, 1599 (1964).
18.
Tseung, A. C. C. & Bevan, H. L. A reversible oxygen electrode. Journal of
ElectroanalyticalChemistry and Interfacial Electrochemistry 45, 429 (1973).
19.
Ruetschi, P. & Delahay, P. Influence of Electrode Material on Oxygen
Overvoltage: A Theoretical Analysis. The Journal of Chemical Physics 23, 556
(1955).
20.
Hoar, T. P. The mechanism of the oxygen electrode. Proceedings of the Royal
Society of London. Series A, Containing Papers of a Mathematical and Physical
Character142, 628 (1933).
21.
Damjanovic, A. Anodic Oxide Films as Barriers to Charge Transfer in 02
Evolution at Pt in Acid Solutions. Journal of The Electrochemical Society 123,
374 (1976).
22.
Damjanovic, A. Electron transfer through thin anodic films in oxygen evolution at
Pt electrodes in alkaline solutions. ElectrochimicaActa 37, 2533 (1992).
23.
Bockris, J. 0. Kinetics of Activation Controlled Consecutive Electrochemical
Reactions: Anodic Evolution of Oxygen. The Journal of Chemical Physics 24, 817
(1956).
24.
Rossmeisl, J., Logadottir, A. & Norskov, J. K. Electrolysis of water on (oxidized)
metal surfaces. Chemical Physics 319, 178 (2005).
25.
Rossmeisl, J., Qu, Z. W., Zhu, H., Kroes, G. J. & Norskov, J. K. Electrolysis of
water on oxide surfaces. Journal of ElectroanalyticalChemistry 607, 83 (2007).
26.
Man, I. C. et al. Universality in Oxygen Evolution Electrocatalysis on Oxide
Surfaces. ChemCatChem 3, 1159 (2011).
27.
Azzam, A. M., Bockris, J. 0., Conway, B. E. & Rosenberg, H. Some aspects of
the measurement of hydrogen overpotential. Transactions of the Faraday Society
46, 918 (1950).
28.
Bockris, J. 0. & Huq, A. K. M. S. The Mechanism of the Electrolytic Evolution of
Oxygen on Platinum. Proceedings of the Royal Society A: Mathematical, Physical
and Engineering Sciences 237, 277 (1956).
62
29.
Burke, L. D. & Moynihan, A. Oxygen electrode reaction. Part 1.-Nature of the
inhibition process. Transactions of the FaradaySociety 67, 3550 (1971).
30.
Visscher, W. & Devanathan, M. Anodic behavior of platinum electrodes in
oxygen-saturated acid solutions. Journal of Electroanalytical Chemistry 8, 127
(1964).
31.
Riddiford, A. Mechanisms for the evolution and ionization of oxygen at platinum
electrodes. Electrochimica Acta 4, 170 (1961).
32.
Hoare, J. P. The Reversible Oxygen Electrode. Nature 211, 703 (1966).
33.
Iwakura, C., Fukuda, K. & Tamura, H. The anodic evolution of oxygen on
platinum oxide electrode in alkaline solutions. Electrochimica Acta 21, 501 (1976).
34.
Appleby, A. Rest potentials of oxide-free platinum electrodes. Journal of
ElectroanalyticalChemistry and Interfacial Electrochemistry 35, 193 (1972).
35.
Hoare, J. P. Oxygen Overvoltage Measurements on Bright Platinum in Acid
Solutions. Journal of The Electrochemical Society 112, 849 (1965).
36.
Rand, D. A. J. & Woods, R. Cyclic voltammetric studies on iridium electrodes in
sulphuric acid solutions. Journal of Electroanalytical Chemistry and Interfacial
Electrochemistry 55, 375 (1974).
37.
Gambardella, A. A., Bjorge, N. S., Alspaugh, V. K. & Murray, R. W.
Voltammetry of Diffusing 2 nm Iridium Oxide Nanoparticles. The Journal of
Physical Chemistry C 115, 21659 (2011).
38.
Buckley, D. N. & Burke, L. D. The oxygen electrode. Part 6. -Oxygen evolution
and corrosion at iridium anodes. Journal of the Chemical Society, Faraday
Transactions 1 72, 2431 (1976).
39.
Burke, L. D. & O'Meara, T. 0. Oxygen electrode reaction. Part 2. -Behaviour at
ruthenium black electrodes. Journal of the Chemical Society, Faraday
Transactions 1 68, 839 (1972).
40.
Beni, G., Schiavone, L. M., Shay, J. L., Dautremont-Smith, W. C. & Schneider,
B. S. Electrocatalytic oxygen evolution on reactively sputtered electrochromic
iridium oxide films. Nature 282, 281 (1979).
41.
Burke, L. D., Murphy, 0. J., O'Neill, J. F. & Venkatesan, S. The oxygen
electrode. Part 8.-Oxygen evolution at ruthenium dioxide anodes. Journal of the
Chemical Society, Faraday Transactions 1 73, 1659 (1977).
42.
Ferrer, J. E. & Victori, L. Study of the oxygen evolution reaction on the iridium
electrode in acid medium by eis. Electrochimica Acta 39, 667 (1994).
63
43.
Ferrer, J. E. & Victori, L. Oxygen evolution reaction on the iridium electrode in
basic medium studied by electrochemical impedance spectroscopy. Electrochimica
acta 39, 581 (1994).
44.
Nakagawa, T., Beasley, C. a. & Murray, R. W. Efficient Electro-Oxidation of
Water near Its Reversible Potential by a Mesoporous IrOx Nanoparticle Film. The
Journal of Physical Chemistry C 113, 12958 (2009).
45.
Owe, L.-E., Tsypkin, M. & Sunde, S. The effect of phosphate on iridium oxide
electrochemistry. Electrochimica Acta 58, 231 (2011).
46.
Owe, L.-E., Tsypkin, M., Wallwork, K., Haverkamp, R. G. & Sunde, S. Iridiumruthenium single phase mixed oxides for oxygen evolution: Composition
dependence of electrocatalytic activity. Electrochimica Acta 70, 158 (2012).
47.
Doyle, R. L. & Lyons, M. E. G. Kinetics and Mechanistic Aspects of the Oxygen
Evolution Reaction at Hydrous Iron Oxide Films in Base. Journal of the
Electrochemical Society 160, H142 (2013).
48.
Lyons, M. E. G. & Brandon, M. P. The oxygen evolution reaction on passive
oxide covered transition metal electrodes in alkaline solution. Part 111-Iron.
InternationalJournal of Electrochemical Science 3, 1463 (2008).
49.
Lyons, M. E. G. & Brandon, M. P. The' oxygen evolution reaction on passive
oxide covered transition metal electrodes in aqueous alkaline solution. Part 1Nickel. InternationalJournal of Electrochemical Science 3, 1386 (2008).
50.
Lyons, M. E. G. & Brandon, M. P. The oxygen evolution reaction on passive
oxide covered transition metal electrodes in alkaline solution. Part 2-Cobalt.
InternationalJournal of Electrochemical Science 3, 1425 (2008).
51.
Burke, L., Lyons, M. & Murphy, 0. Formation of hydrous oxide films on cobalt
under potential cycling conditions. Journal of Electroanalytical Chemistry and
Interfacial Electrochemistry 132, 247 (1982).
52.
Brossard, L. Cobalt black electrodes for the oxygen evolution reaction from
electrolysis of 40 wt% KOH. International Journal of Hydrogen Energy 17, 671
(1992).
53.
Brossard, L. & Messier, C. Nickel coatings: deposition on an iron substrate and 02
evolution in 40 wt.% KOH. Materials Chemistry and Physics 33, 239 (1993).
54.
Jones, E. & Wynne-Jones, W. F. K. The nickel oxide electrode. Part 2.
Transactions of the Faraday Society 52, 1260 (1956).
55.
Conway, B. E. & Bourgault, P. L. The Electrochemical Behavior of the NickelNickel Oxide Electrode: Part I. Kinetics of Self-Discharge. Canadian Journal of
Chemistry 37, 292 (1959).
64
56.
Hickling, A. & Hill, S. Oxygen overvoltage. Part I. The influence of electrode
material, current density, and time in aqueous solution. Discussions of the
Faraday Society 1, 236 (1947).
57.
Risch, M. et al. Water Oxidation by Electrodeposited Cobalt Oxides-Role of
Anions and Redox-Inert Cations in Structure and Function of the Amorphous
Catalyst ChemSusChem 5, 542 (2012).
58.
Risch, M. et al. Atomic structure of cobalt-oxide nanoparticles active in lightdriven catalysis of water oxidation. InternationalJournal of Hydrogen Energy 37,
8878 (2012).
59.
Esswein, A. J., Surendranath, Y., Reece, S. Y. & Nocera, D. G. Highly active
cobalt phosphate and borate based oxygen evolving catalysts operating in neutral
and natural waters. Energy & Environmental Science 4, 499 (2011).
60.
Bediako, D. K. et al. Structure-Activity Correlations in a Nickel-Borate Oxygen
Evolution Catalyst. Journal of the American Chemical Society 134, 6801 (2012).
61.
Bediako, D. K., Surendranath, Y. & Nocera, D. G. Mechanistic Studies of the
Oxygen Evolution Reaction Mediated by a Nickel-Borate Thin Film
Electrocatalyst. Journal of the American Chemical Society 135, 3662 (2013).
62.
Surendranath, Y., Kanan, M. W. & Nocera, D. G. Mechanistic studies of the
oxygen evolution reaction by a cobalt-phosphate catalyst at neutral pH. Journal
of the American Chemical Society 132, 16501 (2010).
63.
Risch, M. et al. Cobalt-oxo core of a water-oxidizing catalyst film. Journal of the
American Chemical Society 131, 6936 (2009).
64.
Risch, M. et al. Characterisation of a water-oxidizing Co-film by XAFS. Journal
of Physics: Conference Series 190, 012167 (2009).
65.
Kanan, M. W. et al. Structure and valency of a cobalt-phosphate water oxidation
catalyst determined by in situ X-ray spectroscopy. Journal of the American
Chemical Society 132, 13692 (2010).
66.
Balandin, A. A. Modern State of the Multiplet Theory of Heterogeneous
Catalysis. Advances in Catalysis 19, 1 (1969).
67.
Sabatier, P. Hydrog6nations et d6shydrog6nations par catalyse. Berichte der
deutschen chemischen Gesellschaft 44, 1984 (1911).
68.
Trasatti, S. Electrocatalysis by oxides-Attempt at a unifying approach. Journal of
Electroanalytical Chemistry 111, 125 (1980).
69.
Koper, M. T. M. Thermodynamic theory of multi-electron transfer reactions:
Implications for electrocatalysis. Journal of Electroanalytical Chemistry 660, 254
(2011).
65
70.
Wolfram, T. & Ellialtioglu, $. Electronic and Optical Properties of d-Band
Perovskites. (Cambridge University Press: New York, 2006).
71.
Peia, M. a. & Fierro, J. L. G. Chemical Structures and Performance of Perovskite
Oxides. Chemical Reviews 101, 1981 (2001).
72.
Bockris, J. 0. M. & Otagawa, T. The electrocatalysis of oxygen evolution on
perovskites. Journal of the Electrochemical Society 131, 290 (1984).
73.
Bockris, J. 0. M. & Otagawa, T. Mechanism of oxygen evolution on perovskites.
The Journal of Physical Chemistry 87, 2960 (1983).
74.
Bockris, J. 0. M., Otagawa, T. & Young, V.- Solid state surface studies of the
electrocatalysis of oxygen evolution on perovskites. Journal of Electroanalytical
Chemistry and Interfacial Electrochemistry 150, 633 (1983).
75.
Otagawa, T. Lanthanum Nickelate as Electrocatalyst: Oxygen Evolution. Journal
of The Electrochemical Society 129, 2391 (1982).
76.
Suntivich, J. et al. Design principles for oxygen-reduction activity on perovskite
oxide catalysts for fuel cells and metal-air batteries. Nature Chemistry 3, 546
(2011).
77.
Tejuca, L. G., Fierro, J. L. G. & Tasc6n, J. M. D. Structure and Reactivity of
Perovskite-Type Oxides. Advances in Catalysis 36, 237 (1989).
78.
Cox, P. A. Transition Metal Oxides: An Introduction to their Electronic Structure
and Properties.(Oxford University Press: New York, 1992).
79.
Matsumoto, Y., Yoneyama, H. & Tamura, H. Electrochemical properties of
lanthanum nickel oxide. Journal of Electroanalytical Chemistry and Interfacial
Electrochemistry 80, 115 (1977).
80.
Matsumoto, Y., Kurimoto, J. & Sato, E. Oxygen evolution on SrFeO3 electrode.
Journal of ElectroanalyticalChemistry 102, 77 (1979).
81.
Matsumoto, Y. Oxygen Evolution on La -Sr Fe ,Co,0
The Electrochemical Society 127, 2360 (1980).
82.
Matsumoto, Y. & Sato, E. Oxygen evolution on La xSr MnO 3 electrodes in
alkaline solutions. Electrochimica Acta 24, 421 (1979).
83.
Matsumoto, Y. Oxygen Evolution on La -SrXCoO 3 Electrodes in Alkaline
Solutions. Journal of The Electrochemical Society 127, 811 (1980).
84.
Hammer, B. Electronic factors determining the reactivity of metal surfaces.
Surface Science 343, 211 (1995).
66
3
Series Oxides. Journal of
85.
Hammer, B. Special Sites at Noble and Late Transition Metal Catalysts. Topics
in Catalysis 37, 3 (2006).
86.
Crumlin, E. J. et al. Surface Strontium Enrichment on Highly Active Perovskites
for Oxygen Electrocatalysis in Solid Oxide Fuel Cells. Energy & Environmental
Science 5, 6081 (2012).
87.
Ferreira, P. J. et al. Instability of Pt/C Electrocatalysts in Proton Exchange
Membrane Fuel Cells. Journal of The Electrochemical Society 152, A2256 (2005).
88.
Suntivich, J., Gasteiger, H. A., Yabuuchi, N. & Shao-Horn, Y. Electrocatalytic
Measurement Methodology of Oxide Catalysts Using a Thin-Film Rotating Disk
Electrode. Journal of the Electrochemical Society 157, B1263 (2010).
89.
Kresse, G. & Hafner, J. Ab initio molecular dynamics for liquid metals. Physical
Review B 47, 558 (1993).
90.
Kresse, G. & Furthmiiller, J. Efficient iterative schemes for ab initio total-energy
calculations using a plane-wave basis set. Physical Review B 54, 11169 (1996).
91.
B16chl, P. E. Projector augmented-wave method. Physical Review B 50, 17953
(1994).
92.
Perdew, J. P. & Wang, Y. Accurate and simple analytic representation of the
electron-gas correlation energy. Physical Review B 45, 13244 (1992).
93.
Anisimov, V. I., Aryasetiawan, F. & Lichtenstein, A. I. First-principles
calculations of the electronic structure and spectra of strongly correlated systems:
the LDA + U method. Journal of Physics: Condensed Matter 9, 767 (1997).
94.
Dudarev, S. L., Savrasov, S. Y., Humphreys, C. J. & Sutton, A. P. Electronenergy-loss spectra and the structural stability of nickel oxide: An LSDA+U
study. Physical Review B 57, 1505 (1998).
95.
Lee, Y.-L., Kleis, J., Rossmeisl, J. & Morgan, D. Ab initio energetics of LaBO 3
(001) (B=Mn, Fe, Co, and Ni) for solid oxide fuel cell cathodes. Physical Review
B 80, 224101 (2009).
96.
Lee, Y.-L., Kleis, J., Rossmeisl, J., Shao-Horn, Y. & Morgan, D. Prediction of
solid oxide fuel cell cathode activity with first-principles descriptors. Energy &
Environmental Science 4, 3966 (2011).
97.
Wei, S.-H., Ferreira, L., Bernard, J. & Zunger, A. Electronic properties of random
alloys: Special quasirandom structures. Physical Review B 42, 9622 (1990).
98.
Hertz, J. et al. Magnetism and structure of LiCoO 2 and comparison to Na.CoO 2.
Physical Review B 77, 1 (2008).
67
99.
Miyazaki, S., Kikkawa, S. & Koizumi, M. Chemical and electrochemical
deintercalations of the layered compounds LiMO 2 (M = Cr, Co) and NaM'02
(M' = Cr, Fe, Co, Ni). Synthetic Metals 6, 211 (1983).
100.
Arnold, M., Gesing, T. M., Martynczuk, J. & Feldhoff, A. Correlation of the
Formation and the Decomposition Process of the BSCF Perovskite at
Intermediate Temperatures. Chemistry of Materials 20, 5851 (2008).
101.
Zhang, W., Wang, Z. & Chen, X. M. Crystal structure evolution and local
symmetry of perovskite solid solution Ba[(Fe,/ 2 Nb1/2 )l _Tix]O 3 investigated
by
Raman spectra. Journal of Applied Physics 110, 064113 (2011).
102.
Ishikawa, a., Nohara, J. & Sugai, S. Raman Study of the Orbital-Phonon
Coupling in LaCoO3 . Physical Review Letters 93, 1 (2004).
103.
Gerken, J. B. et al. Electrochemical water oxidation with cobalt-based
electrocatalysts from pH 0-14: the thermodynamic basis for catalyst structure,
stability, and activity. Journal of the American Chemical Society 133, 14431
(2011).
104.
Risch, M. et al. Nickel-oxido structure of a water-oxidizing catalyst film. Chemical
communications (Cambridge, England) 47, 11912 (2011).
105.
Zaharieva, I. et al. Synthetic manganese-calcium oxides mimic the water-oxidizing
complex of photosynthesis functionally and structurally. Energy & Environmental
Science 4, 2400 (2011).
106.
Grimaud, A. et al. Double Perovskites as a New Family of Highly Active
Catalysts For Oxygen Evolution in Alkaline Solution. Submitted
107.
Risch, M. et al. Structural Changes of Cobalt-Based Perovskites upon Water
Oxidation Investigated by EXAFS. The Journal of Physical Chemistry C 117,
8628 (2013).
108.
McIntosh, S., Vente, J., Haije, W., Blank, D. & Bouwmeester, H. Structure and
oxygen stoichiometry of SrCo08 Fe0 .20 3 _6 and Ba 0 5 Sr 0 5 Co 0 8 Fe0.2 3 -. Solid State
Ionics 177, 1737 (2006).
109.
Burke, L. D. & O'Sullivan, E. J. M. Oxygen gas evolution on hydrous oxides An example of three-dimensional electrocatalysis? Journal of Electroanalytical
Chemistry and Interfacial Electrochemistry 117, 155 (1981).
110.
Buckley, D. N., Burke, L. D. & Mulcahy, J. K. The oxygen electrode. Part 7.
Influence of some electrical and electrolyte variables on the charge capacity of
iridium in the anodic region. Journal of the Chemical Society, Faraday
Transactions 1 72, 1896 (1976).
68
111.
Buckley, D. N. & Burke, L. D. The oxygen electrode. Part 5.-Enhancement of
charge capacity of an iridium surface in the anodic region. Journal of the
Chemical Society, Faraday Transactions 1 71, 1447 (1975).
112.
McIntyre, J. D. E. Oxidation State Changes and Structure of Electrochromic
Iridium Oxide Films. Journal of The ElectrochemicalSociety 127, 1264 (1980).
113.
Wang, L., Merkle, R. & Maier, J. Surface Kinetics and Mechanism of Oxygen
Incorporation Into Ba ,Sr.CoFe -Y0 3- SOFC Microelectrodes. Journal of The
Electrochemical Society 157, B1802 (2010).
114.
of
Kuroda, C., Zheng, K. & Swierczek, K. Characterization
in
IT-SOFC
application
for
perovskites
GdBao.5 Sr 5 Co, Fe,,05InternationalJournal of Hydrogen Energy 8, 1 (2012).
novel
cells.
69